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Solution Stoichiometry

What mass of KI is required to make 500. mL of a 2.80 M KI solution?. moles of solute. M = molarity =. liters of solution. volume KI. moles KI. grams KI. 1 L. 2.80 mol KI. 166 g KI. x. x. x. 1000 mL. 1 L soln. 1 mol KI. Solution Stoichiometry.

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Solution Stoichiometry

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  1. What mass of KI is required to make 500. mL of a 2.80 M KI solution? moles of solute M = molarity = liters of solution volume KI moles KI grams KI 1 L 2.80 mol KI 166 g KI x x x 1000 mL 1 L soln 1 mol KI Solution Stoichiometry The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. M KI M KI 500. mL = 232 g KI

  2. Moles of solute before dilution (i) Moles of solute after dilution (f) = Dilution Add Solvent = MfVf MiVi Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution.

  3. How would you prepare 60.0 mL of 0.200 M HNO3 from a stock solution of 4.00 M HNO3? MfVf 0.200 x 0.0600 Vi = = 4.00 Mi MiVi = MfVf Mi = 4.00 Vi = ? L Mf = 0.200 Vf = 0.0600 L = 0.00300 L = 3.00 mL 3.00 mL of acid + 57.0 mL of water = 60.0 mL of solution

  4. Titrations In a titration a solution of accurately known concentration is gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color

  5. WRITE THE CHEMICAL EQUATION! What volume of a 1.420 M NaOH solution is Required to titrate 25.00 mL of a 4.50 M H2SO4 solution? H2SO4 + 2NaOH 2H2O + Na2SO4 M M rx volume acid moles acid moles base volume base base acid coef. 4.50 mol H2SO4 2 mol NaOH 1000 ml soln x x x 1000 mL soln 1 mol H2SO4 1.420 mol NaOH 25.00 mL = 158 mL

  6. HCl + NH3 NH4Cl What volume of a 0.800 M HCl solution is required to titrate 40.00 mL of a 1.600 M NH3 solution? 1.600 mol NH3 1 mol HCl 1000 ml soln x x x 1000 mL soln 1 mol NH3 0.800 mol HCl = 80.0 mL 40.00 mL

  7. 3Ba(OH)2 + 2H3PO4 6H2O + Ba3(PO4)2 What volume of a 1.200 M Ba(OH)2 solution is required to titrate 90.00 mL of a 3.600 M H3PO4 solution? 3.600 mol H3PO4 3 mol Ba2(OH) 1000 ml soln x x x 1000 mL soln 2 mol H3PO4 1.200 mol Ba2(OH) = 405.0 mL 90.00 mL

  8. Reactions in Aqueous Solution • 71% of Earth’s surface is water • Another 3% is covered with ice • 66% of human body is water Chapter 4

  9. Solution Solvent Solute A solution is a homogenous mixture of 2 or more substances The solvent is the substance present in the larger amount The solute is(are) the substance(s) present in the smaller amount(s) H2O Soft drink (l) Sugar, CO2 Air (g) N2 O2, Ar, CH4 Pb Sn Soft Solder (s)

  10. nonelectrolyte weak electrolyte strong electrolyte An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity.

  11. H2O C6H12O6 (s) C6H12O6 (aq) H2O NaCl (s)Na+ (aq) + Cl- (aq) CH3COOHCH3COO- (aq) + H+ (aq) Electrolyte Conducts electricity in solution? Cations (+) and Anions (-) Strong Electrolyte – 100% dissociation Weak Electrolyte – not completely dissociated Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution

  12. Table 1 Electrolyte Classification of Some Common Substances Strong ElectrolytesWeak ElectrolytesNonelectrolytes HCl, HBr, HI CH3COOH H2O HClO4 HF CH3OH HNO3 C2H5OH H2SO4 C12H22O11(sucrose) KBr Most organic compounds NaCl NaOH, KOH Other soluble ionic compounds E.g. Calculating ion concentrations in a solution of strong electrolyte. What are the concentration (in molarity) of Mg2+ and Cl- ions in 5.0g/L of MgCl2 solution? [Mg2+] = 5.0g x 1 molMgCl2 x 1 mol Mg2+ =5.3 x10-2 mol/L L 95.21gMgCl2 1 molMgCl2 [Cl-] =2 x 5.3x10-2 = 0.11 mol/L

  13. Concentration of Ions A bottle labeled as 0.100 M Al2(SO4)3. [Al3+] = _____ M (mol / L) [SO42–] = _____ M Assume sea water is 0.438 M NaCl, 0.0512 M MgCl2, and 0.001 M CaCl2 [Na+] = _____ M [Mg2+] = _____ M [Ca2+] = _____ M [Cl–] = _____ M Know how to calculate your quantities

  14. Aqueous Reactions Aqueous reactions can be grouped into three general categories, each with its own kind of driving force: 1. Precipitation reactions 2. Acid base neutralization reactions 3. Oxidation-reduction reactions.

  15. PbI2 Precipitation Reactions When ions form a solid that is not very soluble, a solid is formed. Such a phenomenon is called precipitation. Precipitate – insoluble solid that separates from solution

  16. Rules for converting molecular equations to ionic equations • The rules for converting molecular equations to ionic equations follow: • 1) Make sure the molecular equation is balanced • 2) Ionic substances indicated in the molecular equation as dissolved in solution, such as NaCl(aq), are normally written as ions. • 3) Ionic substances that are insoluble (do not dissolve) either as reactants or products (such as precipitate) are represented by formulas of the compounds • 4) Molecular substances that are strong electrolytes, such as strong acids, are written as ions. Thus, HCl(aq) is written as H3O+(aq) + Cl-(aq) or as H+(aq) + Cl-(aq). • 5) Molecular substances that are weak electrolytes or nonelectrolytes are represented by their molecular formulas.

  17. Solubility Rules for Common Ionic Compounds In water at 25oC 2. All ammonium (NH4+) compounds are soluble. 3.All nitrate (NO3-), chlorate (ClO3-), and perchlorate (ClO4-) compounds are soluble. 1. All alkali metals (Group 1A) compounds are soluble. 4. Most hydroxides (OH-) are insoluble. The exceptions are barium hydroxide [Ba(OH)2], which is very soluble, calcium hydroxide [Ca(OH)2], which is slightly soluble, and previous examples. 5. Most compounds containing chlorides (Cl-), bromides (Br-), or iodides (I-), are soluble. The exceptions are those containing Ag+, Hg22+, and Pb2+. 6. Most sulfates (SO42-) are soluble, excepting previous rules. Calcium sulfate [CaSO4] and silver sulfate [Ag2SO4] are slightly soluble. Barium sulfate [BaSO4], mercury (II) sulfate [HgSO4], and lead sulfate [PbSO4] are insoluble. 7. All carbonates (CO32-), phosphates (PO43-), and sulfides (S2-) are insoluble, excepting previous rules.

  18. AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Write the net ionic equation for the reaction of silver nitrate with sodium chloride. Ag+ + NO3- + Na+ + Cl- AgCl (s) + Na+ + NO3- Ag+ + Cl- AgCl (s) Writing Net Ionic Equations • Write the balanced molecular equation. • Write the ionic equation showing the strong electrolytes • Determine precipitate from solubility rules • Cancel the spectator ions on both sides of the ionic equation

  19. write a net ionic equation for the reaction. (a) Al2(SO4)3 + NaOH i) write down the reactants and interchange of anions to get product Al2(SO4)3 + 6NaOH2Al(OH)3 + 3Na2SO4 All common Na compounds are water soluble Na+ remain in solution. The combination of Al3+ and OH- produce insoluble Al(OH)3. Then the ionic equation is 2Al3++3SO42- + 6Na+ + 6OH-2Al(OH)3(s)+ 6Na++ 3SO42- The net ionic equation is : Al3+ + 3OH-Al(OH)3(s)

  20. Net Ionic Equations Note that some ions appear on both side of equation. These ions go through the reaction unchanged- does not take part in the reaction. We called them spectator ions. We can cancel them from the equation. The resulting equation is a net ionic equation. Ag+(aq) + I-(aq) AgI(s) Net ionic Equation A net ionic equation is an equation that includes only the actual participants in a reaction, with each participant denoted by the symbol or formula that best represent it.

  21. precipitate Pb(NO3)2(aq) + 2NaI (aq) PbI2(s) + 2NaNO3(aq) Pb2+ + 2NO3- + 2Na+ + 2I- PbI2 (s) + 2Na+ + 2NO3- Pb2+ + 2I- PbI2 (s) PbI2 Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation Na+ and NO3- are spectator ions net ionic equation

  22. Precipitation Reactions Heterogeneous Reactions Spectator ions Ag+ (aq) + NO3– (aq) + Cs+ (aq) + I– (aq)  AgI (s) + NO3– (aq) + Cs+ (aq) Ag+ (aq) + I– (aq)  AgI (s) (net reaction)or Ag+ + I– AgI (s) Mostly insoluble Silver halidesMetal sulfides, hydroxidescarbonates, phosphates Soluble ions Alkali metals, NH4+nitrates, ClO4-, acetate Mostly soluble ions Halides, sulfates

  23. Acids (according to Arrhenius) • Have a sour taste. Vinegar owes its taste to acetic acid. • Citrus fruits contain citric acid. • Cause certain plant dyes to change color. For example, • they turn litmus from blue to red. • React with certain metals to produce hydrogen gas. • React with carbonates (CO32-) and bicarbonates (HCO3-)to • produce carbon dioxide gas. • Are electrolytes.

  24. Bases (according to Arrhenius) • Have a bitter taste. • Cause certain plant dyes to change color. For example, • they turn litmus from red to blue. • Feel slippery (they make water insoluble organic molecules • into water soluble molecules). Many soaps contain bases. • Are electrolytes.

  25. Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water

  26. A Brønsted acid must contain at least one ionizable proton! A Brønsted acid is a proton donor A Brønsted base is a proton acceptor base acid acid base

  27. HCl H+ + Cl- HNO3 H+ + NO3- CH3COOH H+ + CH3COO- H2SO4 H+ + HSO4- HSO4- H+ + SO42- H3PO4 H+ + H2PO4- H2PO4- H+ + HPO42- HPO42- H+ + PO43- Monoprotic acids Strong electrolyte, strong acid Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid

  28. acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H2O H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O H+ + OH- H2O 2. Neutralization Reaction net ionic equation for strong acid / strong base reaction

  29. acid + base salt + water HCl (aq) + NH3 (aq) NH4Cl(aq) H+ + Cl- + NH3(aq) NH4+ + Cl- H+ + NH3(aq) NH4+ Neutralization Reaction net ionic equation for strong acid / weak base reaction

  30. Oxidation-Reduction Reaction Oxidation is the lose of one or more electrons by a substance. Historically oxidation is combination of an element with oxygen Reduction is the gain of one or more electrons by a substance. Reduction is referred to the removal of oxygen from an oxide A redox reaction is a process in which electrons are transferred between substance or in which atoms change oxidation number. Oxidizing agent: contains an element whose oxidation state decreases in a redox reaction (it make possible for some other substance to be oxidized and itself reduced.) gains electrons (electrons are found on the left side of its half-equation). Reducing agent: contains an element whose oxidation state increases in a redox reaction, loses electrons (electrons are found on the left side of its half-equation).

  31. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H2, O2, P4 = 0 • In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 • The oxidation number of oxygen isusually–2. In H2O2 and O22- it is –1.

  32. Oxidation numbers of all the elements in HCO3- ? • The oxidation number of hydrogen is +1except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. • Group IA metals are +1, IIA metals are +2 and fluorine is always –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4

  33. Fig. 4.10

  34. Oxidation numbers of all the elements in the following ? IF3 F = -1 3x(-1) + ? = 0 I = +3 K2Cr2O7 NaIO3 O = -2 K = +1 O = -2 Na = +1 3x(-2) + 1 + ? = 0 7x(-2) + 2x(+1) + 2x(?) = 0 I = +5 Cr = +6

  35. 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- 2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e- 2Mg + O2 2MgO Oxidation-Reduction Reactions (electron transfer reactions) Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-)

  36. Fe (s) + 2HCl (aq) FeCl2 (aq) + H2 (g) 2H+ + 2e- H2 Fe Fe2+ + 2e- Iron (Fe) reacting with hydrochloric acid (HCl) Fe is the reducing agent Fe is oxidized H+is reduced H+ is the oxidizing agent An oxidizing agent: contains an element whose oxidation state decreases in a redox reaction (it make possible for some other substance to be oxidized and itself reduced.) gains electrons (electrons are found on the left side of its half-equation). A reducing agent: contains an element whose oxidation state increases in a redox reation loses electrons (electrons are found on the left side of its half-equation).

  37. Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s) Cu2+ + 2e- Cu Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu Cu2+ + 2e- Zn Zn2+ + 2e- Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s) Ag+ + 1e- Ag Zn is the reducing agent Zn is oxidized Cu2+is reduced Cu2+ is the oxidizing agent Cuis oxidized Cu is the reducing agent Ag+is reduced Ag+ is the oxidizing agent

  38. Ca2+ + CO32- CaCO3 NH3 + H+ NH4+ Classify the following reactions. Zn + 2HCl ZnCl2 + H2 Ca + F2 CaF2 Precipitation Acid-Base Redox (H2 Displacement) Redox (Combination)

  39. Chapter 4: Reactions in Aqueous Solutions

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